Chapter 9 Liquids and Solids

1. Chapter 9Liquids and Solids

2. Intro Vocabulary • Gas: no definite shape or volume • Remember kinetic theory of gases • Liquid: definite volume – no definite shape • Some attraction between molecules or atoms • Solid: definite shape and volume • Strong intermolecular bonding 1. Molecules are much closer together in liquids and solids than in gases • In gases, molecules are separated by ten or more molecular diameters • In liquids and solids, the molecules are in contact with each other 2. Intermolecular forces play a major role in the behavior of liquids and solids, whereas they are negligible in gases

3. B. Phase Transitions • Q = m c ∆T (use for temperature changes) • Q = n ∆H (use for phase changes) • Temperature DOES NOT change during a phase change

4. Phase transitions • Melting/Freezing • Heat of fusion – energy required to melt/freeze 1 mole of a substance • Vaporization/Condensation • Heat of vaporization – energy required to vaporize/condense 1 mole of a substance • Sublimation/Deposition

5. Heating/Cooling Curve

6. Phase Diagrams • A. Heating/Cooling Curves definitions • 1. Conversion of a solid to a liquid is:_______________ • 2. Conversion of a liquid to a solid is:_______________ • 3. The freezing point = melting point • 4. Energy needed to melt a given quantity of solid is • called the ___________________________________.

7. Examples: • Example: How much energy is required to melt 100.0 grams of ice? The heat of fusion is 6.01 kJ/mole.

8. Examples: • Example: How much energy in kJ is required to heat 100.0 grams of liquid water from zero to 100°C, and then vaporize all of it? ∆Hvap= 40.79 kJ/mole

9. 9.1 Liquid - Vapor Equilibrium A. Vaporization (evaporation) process in an open container - evaporation will continue until all the liquid is gone - the energy required for vaporization comes from the surroundings and system - vaporization leaves the remaining liquid cooler - evaporation will occur below the boiling point of a substance - evaporation below the boiling point is slower than at the boiling point

10. B. Enthalpy of vaporization • Definition – the amount of energy change that occurs during the vaporization of 1 mole of a substance q = n ∆Hvap

11. C. Vapor Pressure – the pressure of the gas above a liquid in a closed container; dependent on temperature 1. Closed container vs. open container • In an open container the system includes the surroundings and the liquid will evaporate • In a closed system the liquid will evaporate and begin to condense when equilibrium is established between the liquid and gas

12. 2. Dynamic Equilibrium • When the rate at which the liquid vaporizes is equal to the rate at which the vapor condenses • The liquid level in the container does not change • Molecules are constantly moving between phases with no net change

13. 3. Pressure and Volume • As long as some liquid remains when equilibrium is established, the equilibrium vapor pressure will be the same regardless of the volume of the container

14. E. Vapor Pressure Curves and Temperature 1. Relationships • Vapor pressure of liquid increases as temperature increases

15. 2. General Graph • What does this graph tell you about the relative attraction between molecules for substances a-e?

16. Boiling Point 1. Definition: • The boiling point is the temperature at which the vapor pressure equals atmospheric pressure 2. Normal Boiling Point: • The boiling point at exactly 1 atm of pressure 3. Dependency on pressure • At a certain temperature, large bubbles form throughout the liquid; i.e., the liquid boils • The temperature at which a liquid boils depends on the pressure above it

17. Dependency on pressure (continued) • At high elevation, atmospheric pressure is lower, so the boiling point is lower • To elevate the boiling point and allow food to cook more quickly, a pressure cooker can be used

18. 9.3 Intermolecular Forces • Molecules are the structural units of covalently bonded compounds • Properties of molecules: • nonconductors of electricity when pure • insoluble in water but soluble in nonpolar solvents • low melting points • These properties depend on the intermolecular forces between the molecules

19. A. 3 types of intermolecular forces • Dispersion • Dipole-dipole • Hydrogen bonding

20. Dispersion Forces • Definition – a force of attraction between molecules that is caused by temporary dipoles dipole – a molecule with a positive a negative end

21. Dispersion Forces 2. All molecules have some dispersion forces acting between them temporary dipoles form as a result of the natural movement of e-s in the e- cloud creating areas of positive and negative charge

22. Dispersion Forces 3. Strength of dispersion forces - all molecules have dispersion forces - strength increases with increasing # of e-s

23. Dispersion Forces 4. Dispersion forces increase as molar mass increases - directly proportional - Why? As molar mass increases the # of e-s increases - higher Dispersion forces = higher boiling and melting points because molecules tend to “stick together”

24. Example 9.3 Account for the fact that chlorine is a gas, bromine is a volatile liquid, and iodine is a volatile solid at room temperature.

25. C. Dipole-Dipole Forces • Definition and example - a force of attraction between molecules that is caused by permanent dipoles - CO = polar bond resulting in permanent dipole

26. 2. Higher bp and mp than expected because of the D-D forces the molecules “stick together” and require much more energy for the phase change

27. Figure 9.9

28. Example 9.4

29. D. Hydrogen Bonds 1.Unusually strong type of dipole force • H attached to an N or O or F • The H from one molecule is strongly attracted to the negative end of the dipole of another • The strong dipole forms from the large difference in electronegativities of H and (N, O, or F) 2.Hydrogen bonds are the strongest intermolecular force • unusually high boiling points (H2O vs. CH4) • Small size of H allows the unshared pair from the negative end of the dipole to approach the H closely

30. 4. examples of hydrogen bonding • H2O, NH3, HF

31. 5. Unusual properties of Water • Because of H-bonding: • High specific heat • High boiling point • Liquid phase more dense than solid phase = ice floats

32. Figure 9.10

33. Example 9.6

34. What types of intermolecular forces are present in the following substances? Rank these substances in order of increasing bp. N2 HF SiCl4 CH3Cl NH3

35. Covalent vs. Intermolecular Forces • Three types of intermolecular force • Dispersion • Dipole • Hydrogen bond • All three intermolecular forces are weak relative to the strength of a covalent bond • Attractive energy in ice is 50 kJ/mol • Covalent bond in water is 928 kJ/mol